Periodic Table
Periodic Table
Periodic Trends
Atomic
Size
Ionic
Size
Ionization
Energy
Electronegativity
Origin of The Periodic Table:
Early 1800's Johann Dobereiner – Triads -The properties of middle element are the average of
the other two.
John Newlands - Law of Octave – every first and eighth element have similar
properties.
fits the first 14 elements – after Lithium.
1869 Dimitri Mandeleev
Henry Moseley
- Russian Chemist - arranged elements by increasing atomic
mass.
- 1st to publish a classification system based upon the periodic
law
- Periodic Table
- Discovered that properties were related to atomic number and
not atomic mass.
- Rearranged the Periodic Table to the way we know it today.
Modern Periodic Law – The physical and chemical properties of the elements are a periodic
function of their atomic number – electron configuration
Reading the Periodic Table:
Periods
- horizontal rows - numbered 1 through 7 .
- Period number is equivalent to the outermost principle energy level containing
electrons - valence shell.
Groups or Families -
Vertical columns ( 18 columns).
Numbered 1 through 18.
Old numbering system used Roman numerals and an "A" or "B".
Elements of the same group exhibit similar properties.
Note: 1. The Group number is equivalent to the number of valence electrons.
2. Electron arrangements of elements in a group are similar with the same
valence shell configuration.
Names of Groups and Periods:
Group # 1(1A) - Alkali Metals: - most reactive metals and therefore occur in nature only in
compounds.
- Reduced to free state by electrolysis of their melted (Fused)
compounds.
- reactivity increases with increasing atomic number.
Ex: Na is more reactive than Li
Group # 2(2A)- Alkaline Earth Metals: not as reactive as group 1 but also occur in nature only in
compounds and reduced by electrolysis of their fused compounds.
- reactivity also increases with increasing atomic number.
Ex: Ca is more reactive than Mg but less reactive than K.
Groups # 3-11 - Transition Metals: (Heavy metals)
- incomplete d sublevel.
- ions are usually colored, both in solid compounds and in solution.
- Fe and Co – Ferromagnetic – highly attracted to magnet.
Group # 13 (3A) - Boron Family:
-
- Form 3+ (B, Al, Ga, In) or 1+ (Ga, In, Tl) ions
B – semi metal, Rest are metals
B found in Borax – cleaning agent.
Al – used to make cans and foils.
Group #14 (4A) – Carbon Family: - C -Carbon; most important non-metal. Si and Ge – semimetals,
Sn (Tin) and Pb (Lead) - metals
- Carbon appears in 4 allotropes – Diamond, Graphite, coke, and
Flurofan (buckyball)
- Si combines with oxygen to form silica – found in quart, sand,
and glass. Used as a semi-conductor in computer industry.
Tin – used to make bronze for decorative items.
Lead – use in pipes in past. Toxic. Used in car batteries
Group # 15(5A)- Nitrogen Family: - 1st two elements are nonmetal, next two elements are
semimetals.
- Nitrogen is relatively inactive at STP due to triple bond.
- Nitrogen compounds are relatively unstable.
- Nitrogen and Phosphorous compounds are essential
constituents of all living matter.
- Nitrogen used to make NH3, and nitric acid – used in
fertilizers, dyes, and explosives.
Group # 16 (6A) - Oxygen Family: - Form ions with a –2 charge.
- Show progression from nonmetallic to metallic properties.
- Oxygen is a highly active nonmetal. Diatomic molecule.
- Oxygen may form compounds with most elements.
- Sulfur – has 10 allotropes, used to make sulfuric acid.
- Se and Te are rare elements. - semimetals
- Po is a radioactive element and has no stable isotopes.
Group 17 (7A) - Halogens: - Nonmetals that are very active. Occur as diatomic molecules
- do not occur free in nature.
- activity increase with decreasing atomic number - F is most active.
- Occur as F2(g) , Cl2(g) , Br2(l) , I2(s)
Note: F2 , Cl2 , Br2 , I2 are Diatomic Elements which consist of 2 atoms
joined - found in nature this way. (Others ore H2, N2, and O2, )
- F – Used in toothpastes.
- Cl – used as disinfectant (In bleach).
- Silver compounds of I and Br (Only liquid non-metal) (AgI and AgBr)
coat photographic films.
- I – Important for regulation of thyroid gland.
Group 18 (8A)- Noble gases: - Known in the past as inert gases.
- monoatomic gases (molecules with 1 atom).
- not known to form compound with other elements.
- Xenon and Krypton have formed compounds in the lab with fluorine.
- have completely filled valance shell.
- have stable electron configuration - 8 valence electrons.
- He : Used to fill balloons and Diving tanks
- Ne : with other gases, Used to form light displays
- Ar – most abundant noble gas – provide inert atmosphere
- Ar and Kr – protect filaments in light bulbs.
Note: Reactivity of nonmetals decreases with increasing atomic number.
Periods (Horizontal Rows):
Period # 6 Inner Transition Metals - the Lanthanides - Elements 58 - 71.
Period # 7 Inner Transition Metals - the Actinides - Elements 90 - 103. Include the
transuranium elements – synthetic, radioactive elements with atomic # higher than 92.
Blocks:
Block s - Groups 1 and 2 – added electrons fill-up the s sublevel.
Block p- Groups 13-18 – added electrons fill-up the p sublevel.
Block d - Groups 3-12 – added electrons fill-up the d sublevel prior to the valence shell. Include
transition elements.
Block f - Inner Transition Elements - Lanthanide and Actinide series - electrons fill-up the
f sublevel of 2 levels prior to the valence shell (4f – Lanthanides, 5f – Actinides)
Metals, Nonmetals, and Metalloids (Semimetals):
Elements to the left of the dark zigzag line are metals. More than 2/3 of elements are metals.
Metals physical properties are: Solid at room temperature (Mercury an exception), Luster,
Malleability, Ductility, Good heat and electrical conductivity.
Elements touching the zigzag line are metalloids - have properties of both metals and nonmetals.
Elements to the right of the zigzag line + Hydrogen are nonmetals. Nonmetals physical properties:
Generally gases or brittle dull solids at room temperature (Bromine an exception) poor conductors
of heat and electricity.
In the same period, elements in Group 1 are more active than elements in Group 2.
Atomic Radius (Reference Table S):
Half the distance between adjacent (nearby) nuclei. Three types of atomic radius:
1. Covalent atomic radius - Half the distance between covalently bonded nuclei of the same atom.
2. Van der Waals radius - half the distance between nuclei of the same atom that have not form a
bond.
3. Atomic radius in metals - Half the distance from nucleus to nucleus in a crystalline metal.
Trends in Atomic Radius:
As you go across a period from left to right, the radius decreases due to an increase in nuclear
charge.
As you go down a group, the atomic radius of each successive (following) element is bigger due to
an extra shell.
Ionic Radius:
When a metal atom loses an electron, it becomes a positive ion and the radius gets smaller.
When a nonmetal atom gains an electron, it becomes a negative ion and the radius gets larger.
Ionization Energy, Electron Affinity, and Electronegativity (Reference Table S):
Ionization energy is the amount of energy needed to remove an electron. It is a reflection of how
strongly an atom holds its outermost electron. The smaller the amount of ionization energy the
easier it is to lose an electron.
Electron Affinity is the measure of an atom’s attraction for an extra electron. It represent the
energy change that occurs when an atom gains an extra electron. The smaller the amount of
ionization energy, the easier it is to add an electron.
Electronegativity is the attraction for electrons. The larger the electronegativity, the more the
atom attracts electrons.
Metals have low ionization energy, low electron affinity values, and low electronegativity.
Nonmetals have high ionization energy, lower electron affinity values, and high electronegativity.
Noble gases have the highest ionization energy, high electron affinity values, and no
electronegativity values.
As you go across a period - the ionization energy and electronegativity increases.
As you move down a group - the ionization energy and electronegativity decreases.
Electron affinity has no specific pattern across a period or down a group.
Chemical Reactivity:
In Metals: chemical reactivity increases as you move down a group and decreases as you move along
a period from left to right.
Most active metal: Bottom left hand corner - Fr (francium)
In nonmetals – chemical reactivity decreases as you move down the group, and increases as you
move across a period from left to right.
Most active nonmetal: Upper right hand corner - F (fluorine).